A. L. Crumbliss and M. E. McCabe Ill Duke University Durham. North Carolina 27706 J. A. Dilts and Harvey 6. Herman University o i ~ o r t hCarolina at Greensboro Greensboro. 27412
An Introductory Chemistry Synthesis and Kinetics Experiment
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Utilizing a low cost linear-log output spectrophotometer modification
We present here an inorganic oxidation-reduction kinetics experiment suitable for use in an introductory laboratory program. In addition, we describe a low cost (less than $50) modification of the Bausch and Lomb Spectronic 20 colorimeter which will provide a linear output in absorbance. This output signal may then he fed into an inexpensive strip chart recorder to obtain absorbance values as a function of time. The reaction used for the kinetic study was first reoorted by Haim.'
+
-
H+
CO(NH&N~~+Fe(H20)a2+
Co(HzO)eZ++ 5NH4++ Ng-
+ Fe(Hs0)63+ (1)
field theory. Of equal importance is that the student obtains a feeling for the importance (and satisfaction) of fully characterizing a substrate before starting a kinetic study. The kinetic data for the student exoeriment was collected on a Bausch and Lomh Spectronic 50 Spectrophotometer, modified as described below. These data are consistent with that obtained using a Beckman Acta 111 double beam recording spectrophotometer. The reaction was monitored a t 525 nm where the only species which has an appreciable absorbance is the C O ( N H ~ ) ~ ion N ~(c~=+ 240 cm-'M-'1. Under the conditions described below for the student experiment, the total change in measured absorbance is 0.8 absorbance units. Figure I 1s a graph of pscudo fmt-order rateronstants for reacrmn (1) at 2S0C riron\lll inexcess) plotted a s a f u n c tion of iron(1I) concentration." linear least-squares analysis as the slope, which of these data gives 8.0 X 10-'M-'min-' corresnonds to the second-order rate constant for reaction (1) a t 25"'C. This is in reasonable agreement with the publishkd value of 5.2 X 10-'M-'min-' . Th'IS dlscrenancv ' is a5 exnected ~~r - - ~ - according to our detailed determination of the influence of ionic strength and specific anion effects in this reaction."The published study of Haim was performed under controlled conditions excluding air and a t a constant oerchlorate ion c o n c ~ n t r a t i o ni.nikr .~ the conditinns of the student experiment described here, air uxidatinn of ihe stork irontI11 . . solution is not significant and special precautions to exclude air are not necessary. Ionic strength for the student experiment varies from 0.5-0.8 M. This, however, does not cause any serious difficulties as is evidenced by the linearity of the plot shown in Figure 1. The student kinetic experiment as described below can be performed a t various temperatures. Data collected over a temperature range from 25-4O0C result in a calculated activation enthalpy of 14 i 2 kcal/mole.
While oxidation-reduction reactions are discussed in the lecture portion of introductory courses, they are often overlooked when develooing laboratow kinetics exoeriments. A partic~llaradvantag; ut:thissyatem is that tourpseudo firstorder kinetic experiments at different excess ir(m\ll) concentrations can be completed in a normal 4 hr laboratory period to obtain a secood-order rate constant for reaction (1). Variable temperature studies may be added as an additional experiment to obtain an activation enthalpy value. In addition, the student gains experience in the synthesis of a Werner type inorganic complex which can also be prepared in a normal 4-hr lahoratorv oeriod. Suoolementarv related exneriments may also be used to give the iaboratory program an integrated approach. These include physical characterization and analysis of the substrate complex by EDTA titration to obtain the percent cobalt content, determination of charge on the complex by ioo-exchange, and determination of NH3 c ~ n t e n t . ~ Each experiment can be easily performed in a single laboratory period and provides the student with experience in quantitative pipetting and titration techniques. An additional spectroscopic experiment may be included to determine the infrared and visible absorption spectra of [ C O ( N H ~ ) ~ N $ C I ~ Description of Eleclronics This allows for some laboratory discussion of shifts in infrared Use of a logarithmic transducer with a Spectronic 20 has absorptions upon coordination of azide ion and simple crystal been described previously by Habig, et. aL5 T h e circuit described here is based on modern integrated circuit operational amplifiers and can be constructed for less than $50. With this logarithmic amplifier, it is possible to convert any Spectronic
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A
' Haim, A,, J.Amer. Chem. Soc., 85,1016(1963).
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528 / Journal of Chemical Education
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%Thetechnique used for ammonia analysis is distillation of a weighed sample from strongly basic solution, collecting the distillate in standardized acid and back titrating. See for example, Sebera, D. K., J. CHEM. EDUC., 40,476(1963);Vogel, A. I., "A Textbook of Quantitative Inorganic Analysis," Third Ed., Longmans, London, 1962, p. 254. Figure 1 illustrates an additional advantage in choosing the szidopentaaminecobalt(II1) complex for this student experiment. Aquation of the substrate complexas a camplicatingsidereaction is insignificantin this system, as evidenced by the f a d that the straight line passes through the origin. Crumhliss, A. L., and M~Cabe111, M. E., unpublished results. Readers interested in experiments which demonstrate the influence of ionic strength on the second-order rate constants for a reaction between two +2 ions for use in an advanced laboratarydealing with electrolyte solutions or kinetics should contact the authors. Hahig, R. L., Pardue, H. L., and Worthington, J. B., Anal. Chem., 39,600(1967).
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Fogure 1 Pseao first-xoer rate m s m t . k'. as a fnction01 lranl1)mcemation for react on ( I . Cond lions. ICO~N~,I,N,~'I = 3.5 X lO-'M: Id'] = 0.40 M: I = 0 51 - 081 M, YICI-I = 0.47 - 0.67 M; T = 25'C.
~
Figure 2. Circuit diagram showi- transistor in feedback loop ol operational amelilisr.
20 to a fixed wavelength recording spectrometer with linear presentation of absorbance. It is possible to generate the logarithm of a current or voltage signal by placing a transistor in the feedback loop of a n operational amplifier as shown in Figure 2. The output of such a device can be shown to he approximated by eqn. (2) -Eo = aTlog (ii,)
- bT
(2)
where ii, is the collector current, T is the absolute temperature, and a and b are constants. From this eauation it is seen that there are two temperature effects for which compensation is needed for accurate logarithmic generation. The first efrect is on the slope of the-output which increases by 0.1985 mvPK and the second involves the intercept of the output which changes by about 2 mvIoK for most silicon transistors. For precise logarithmic reaction, both temperature effects should be cancelled out. In the loe-ratio amolifier shown in Figure 3, coefficient b is eliminatedLy use of akatio circuit and emolovine a dual transistor. The effect of temperature on the slope is eyiminated by the resistor network associated with IC4. Modification of the Spectronic 20 consists of installation of a phone jack. When the plug is inserted into the jack, pin 8 of the phototube is connected to the logarithmic amplifier and disconnected from the amplifier in the Spectronic 20 (see Fig. 3). When the plug is removed, the phototube is connected to the Soectronic 20's amnlifier and the instrument can be used in k s normal mode oE operation. The current from the nhototube is initiallv amnlified and ~ converted to a voltage siA;,nal by I C I W H f of a dual transistor, (31,' is placed in the feedback lmos of IC'2 and 1C3 to eenerate the logarithm of the voltages supplied to their inputs. Ijse of the dual transistor aids in temperature compensation. Both IC2 and IC3 are 741 operational amplifier^.^ The outputs of IC2 and IC3 are then fed to IC4 (a 741 o~erationalamnlifier) whose output is a measure of the log r&io of the two input currents. The effect of temperature on the slooe is eliminated by the resistor network associated with 1e4. Rt is a 10K thermistor9 mounted in contact with the dual transistor, Q l . The values of RI and R2 (10K trimpot resistors) determine the desiredmvldecade output of the circuit. In the case where full scale (250 mv) corresponds to 1 absorbance unit, they have n o d a l values of 3K for RI and 2.5K for R2. As described below. the 10 turn 1K ootentiometer associated with the innnt of IC3 is used to set zkro absorbance. Operation of the modified Spectronic 20 connected to a strip chart recorder is as follows. In the conventional fashion without any connection to the loearithmic amnlifier. the Spectronic 20 is set to read -A and zero A with th;light beam blocked and with a reference solution (usuallv distilled water) in the light beam, respectively. A sdution whose ahwnbance can be cmveniently rrad on the meter of the Soectronic 20 is inserted into the i;strument and the absorbanie determined. The distilled water reference solution is then reinserted and the logarithmic amplifier connected to the Spedronic 20. The ten-turn potentiometer is used to zero the recorder, which has previously been zeroed by shorting the input terminals. T h e needle on the meter of the Spectronic 20 will be driven off scale ~
~~
--
Figure 3. Circuit diagram far log-ratio amplifier. when the logarithmic amplifier is connected and may be returned to any position on scale by use of the set zero percent transmission control on the instrument. The light control should not be adiusted. With the solution of known ahsorhante in the lightpath, the appropriate adjustment is made with the recorder span adjustment to give the corresponding absorbance reading on the chart. T h e recorder used in this work was a Heath model EU 208 operating on the 250 mv range; other recorders having a span adjustment should he satisfactory. ~
Student Experiment Synthesis of Azidopentaamminecoban(lllj C h l ~ r i d e ' ~ I ~o a 5UO ml-Erlmmeverflask containing 14Umluf H?O.stepuiic
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add 26gof CKI~GHIO. 7 . d of ~ KnK .2Ogot NH,CI, and 1111ml u i
NH,OH. Stir untilall reacwnrr hnverlim,lved. Warm to-70°C and vigorously bubble air through the solution for 30 min. Cool the solution on an ice-water bath for 10 min and then add 60 ml of fi M HCI. Allow the solution to remain on the ice bath for an additional 20 min to ensure maximum precipitation of the product. Filter the product and wash three times with 30ml of 95%ethanol followed by two washings with 30-ml portions of ether. Air dry while still in the filter funnel to give a deep purple microcrystalline solid. The product may he recrystallized by dissolving in 1 volume of water at room temperature (minimumamount) with the subsequent addition of 112 volume of 6 M HCI and 1 volume of 95% ethanol. Cool the solution and filter; wash the product as indicated above. roncd
Cobalt Analysis by EDTA Titration
It is necessary that all cobalt present have an oxidation state of +2 and that no complering ligands be present to compete with EDTA or murexide. These conditions are best met bv heating the eomoound in either concentrated nitric or sulfuricacidas follo&. weigh' three clean dry porcelain dishes. To each add 20-40 mg of the compound and reweigh the dishes to obtain the weight of sample in each. Add 5.0-7.0 ml of eoncd acid to each dish, cover the dish with a watch glass, and place on a hot plate set at its highest temperature. When colored gas is no longer produced and the resulting solution is clear with a light pink tint, remove the dishes from the heat and allow to cool. It is important that the dishes be removed as soon as these conditions are met to prevent excess boiling or splattering which will lead to loss of the compound. Murexide, the ammonium salt of purpuric acid, is used as the endpoint indicator in the EDTA titration. A sample is prepared by mixing murexide with KC1 in the ratio 1:100. Each titration will require 0.05-0.08 g of this mixture. Acceptable results may he obtained by drying disadium dihydrogen ethylenediaminetetraacetate (EDTA) at 80.0°C for 3 hr and preparing a stock 0.01 M solution by dissolving 3.77 gin 11of HzO. If desired, " I n Intersil uprralional amplifier modrl MUuX:, $6 in unit quan5 h~berElecrronicr, 1u Sunnybrook lirive, Rslelgh, N.C. 2'liltJ "National, 2N3806, $0.30 ammoi id Electronics, 2923 Pacific, Greensboro, N.C. 8Ap~rovimatecost $1 each, Teehnieo Electronic Supply, 6223 Triport Court, Greensboro, N.C. For example, Victory 41D2, $0.88. lo Adapted from "Gmelins Handbuch Der Anorgonixhen Chemie," Vol. 58, Part B, Verlag Chemie, 1964, p. 407.
titit.;.
Volume 53, Number 8, August 1976 / 529
further standardization may be achieved by titration with MgSO4." Toperform the EDTA titration far cobalt, carefully transfer the solution from each porcelain dish into a labelled 300-ml beaker. Be sure to rinse each dish several times with water to ensure complete transfer of all the cohalt(I1) present. Bring the total volume to 150ml and add 6.0 M NHIOH to raise the pH to 8.0-9.0. The pH is a crucial factor in this titration, since at both very high and very low pH murexide will decompose. A pH meter can be used, but if one is not available, add NH40H just until the solution turns a light yellow. Cobalt(I1) in the presence of NHlOH will turn yellow at pH 8.5-10.5; however, addition of excess NHlOH will precipitate cobalt hydroxide. Next add the indicator. The solution will now have a noticeably darker yellow color. Begin titrating the solution with EDTA. Occasionally the solution will begin to turn orange due to protonation of the murexide by EDTA; however, if a few ml of 6 M NH40H are added the color will return to yellow. When theaddition of NH40H nolonger affectsthe color of the solution, the endpoint is near. Continue titrating until no trace of orange is present and the solution is a deep violet color. Each mole of cobalt present requires one mole of EDTA at the endpoint. Use this relationship to calculate the moles of cobalt(I1) present in each sample. The known weight of the sample can then be used to calculate the percent cobalt by weight in the sample. Kinetics of the Reduction of CC(NH~)~N~~+ by Fez+ (aq)
Cobalt normally exists in oxidation states of +2 or +3. Comolexes in which cobalt has a n oxidation state of +3 tend t o beLfairlystable in aqueous solutions, but can be reduced by other aqueous transition metals such a s Fe(I1). This experiment dials with a redox reaction of this type involving C O ( N H ~ ) ~ and N~~ Fe(H20)e2+. + T h e reaction proceeds via a binuclear azido bridged complex with the free end of t h e azide replacing one of the Fe(Hz0k2+ waters during water exchange, which is very rapid for Fe(H20)G2+. Once the bridge has formed, an electron can be transferred from t h e iron throueh the a i d e to the cobalt. T h e cobalt c o m ~ l e xthen exists in t h e reduced labile form, and undergoes rapid ligand exchange to a n eauilibrium state of Co(H?OL2+.T h e rate of the over& reactiodcan be studied by m o n i k n g spectral changes a s the C O ( N H ~ ) ~ is N reduced ~ ~ + and Co(H20)e2+ forms.
Vogel, Ref. (2) Chap. IV. '2A. L C. wishes to thank Professor W. F. Gutknecht for helpful discussions and the Research Corporation for a Cottrell Grant used to purchase the Beekman Acta I11 recording speetrophotometer used in the early development stages of the student kinetics experiment. 1'
530 / Journal of Chemical Educat7on
T h e solutions to b e prepared and the procedures for obtaining the necessary data are a s follows. A) Apparatus a) Visible spectrophotometer h) Water bath at -60°C c) 4-spectrophotometer cuvets-dry B) ~olutions a) 100 ml solution: 0.070 M FeCh and 0.40 M HCI h) 100 ml solution: 0.140M FeCh and 0.40M HCI c/ 100 ml solution: 0.210 M F e ~ l n a n d 0.40 M HCI dl 100 ml solurivn: 0.2A0 A l h'eCL and 0.40Af HCI e) 100 ml rolutron: 7.0 X 10-.'.4 [Co(NH,l:,N1]CI, and 0 40M HCI C) Procedure a) Allow all solutions to come to room temperature. hl Set the ~pertmphotorneterat 225 nm. rl Perform the fc,llowmg for each iron solution la) through (d) 11 Pmet 10 rnl of the FeCb solutim and 10 rnl of the cobalt soiution into a 25-ml beaker and stir for 1 min, 21 Fdl a cuvet with this reaction solution and place it in the spertrophotomctcr. Tnkcnhsorhnnee rendinp for the above solutionio\,er rhe follou,ingperiod of time: n~ fill mi": bj 30 min; c) 20 min; d) 15 min. d) After all four reactions have been studied, place all four cuvets in water at -60°C for 20-25 min and record the absorbance of each at 525 nm. T h e rate of the electron transfer reaction studied can b e expressed as follows
If [Fez+] >> [CO(NH&N~~+], then [Fez+] will remain constant during the reaction and a modified equation can be written
where k' = k[Fe2+] and k' is called a pseudo first-order rate constant. Equation (3) can be integrated t o give the logarithmic relationship = k't I ~ [ C O ( N H ~ ) ~ N~ ~ln[Co(NHd~N3~+] +]~-O
Since absorbance is proportional t o [ C O ( N H & N ~ ~ +a ]t 525 nm, ahsorhances can be substituted into the above equation, and after some rearranging one obtains - (At -At=-) = k't ( A M - At=,) Thus, a plot of -In (At - At,=)l(At=o - A t = , ) a s a function e T h e k' obtained for o f t will vield a straieht line of s l o ~ k'. each [ ~ e $ +can ] then ibe plotted as sfunction of Fez+] and the second-order rate constant k obtained as the s l o ~ eSuch . a olot for iron(I1) solutions a,b,c and d is shown in Figure 1.12
